AP Chem Unit 3

Intramolecular Forces

Intramolecular Forces

• occur between elements inside of a compound

• Ionic, Covalent, Network Covalent, and Metallic

• Stronger than intermolecular forces

Intermolecular Forces

occurs between covalent molecules (IMFs)

London Dispersion Forces

• Present in non-polar covalent, polar covalent, and non-metals (all molecular covalent covalent compounds)

• Weakest IMF

• Low melting and boiling points

• Electron cloud density gets momentarily distorted, creating a very temporary dipole moment (POLARIZABILITY), so as the # of electrons increases, the chance of polarizability increases

  • “More total electrons leads to a more polarizable electron cloud, which causes increases temporary dipole-moments, which leads to stronger london dispersion forces and stronger attractions.”

• polarizability can cause a non-polar covalent substance to have a higher boiling/melting point than a polar covalent substance

• if two compounds have the exact same elements and # of electrons, then the compound with the most surface area will be the post polarizable and have stronger London Dispersion Forces

Dipole-Dipole

• present in polar covalent molecules

• partial positive and partial negative ends always create attractions with other molecules

• medium force

• medium melting and boiling points

• the greater the dipole-moment, the stronger the dipole-dipole forces

Hydrogen Bonding Forces

• present in polar covalent molecules with H-F, H-O, or H-N bonds

• particle positive and partial negative always attract

• the large difference n electronegativity between H and F,O, or N creates a strong dipole moment

• the more hydrogen bonds, the stronger the attractions

• strong force

• high melting and boiling points

Volatility

the ease of evaporation

strong IMFs cause _ volatility and weak IMFs cause _ volatility

low, high

Vapor Pressure

• the pressure exerted by gaseous molecules at the surface of a liquid

• always present in a liquid no matter the temperature since some molecules will always have enough energy to overcome IMFs

At a fixed temperature, strong IMFs cause _ vapor pressure and weak IMFs cause _ vapor pressure

low, high

Increasing temperature causes an _ in vapor pressure

increase

Decreasing temperature causes an _ in vapor pressure

decrease

Vapor pressure “over water” =

total pressure - vapor pressure

Boiling occurs when …

vapor pressure = atmospheric pressure

Phase equilibrium

when the rate of vaporization and condensation in a closed container are equal

Answering Free Response questions over IMFs

1. Identify all the attractions/present in both substances

2. state which substance has greater/weaker attractions

   1. Different forces? stronger force = greater attractions

   2. If both have H-bonding, more H-bonding = greater attractions

   3. If both have the same forces and # of electrons, more surface area = greater attractions

   4. If results are unexpected, it’s probably due to strong London Dispersion Forces (due to polarizability)

3. Relate strength of attractions to the property, and make sure to answer the question asked

If you’re asked to explain why SiO₂ has a higher melting/boiling point than another covalent molecule,

it’s most likely because SIO₂ is Network Covalent and the other molecules is molecular covalent!!

STP conditions

273 K/0° c and 1 atm

Ideal Gas Law

• PV = nRT

• PV = (m/M) RT

• M = DRT/P

Combined Gas Law

P₁V₁/n₁T₁ = P₂V₂/n₂T₂

Boyle’s Law

• P₁V₁ = P₂V₂

• ↑ P = ↓ V

• ↓ P = ↑ V

Charles’ Law

• V₁/T₁ = V₂/T₂

• ↑ T = ↑ V

• ↓ T = ↓ V

Gay Lussac’s Law

• P₁/T₁ = P₂/T₂

• ↑ T = ↑ P

• ↓ T = ↓ P

Avogadro’s Law

• V₁/n₁ = V₂/n₂

• ↑ n = ↑ V

• ↓ n = ↓ V

• equal volumes of gases at the same temperature and pressure contain equal #s of particles (moles) even if the gases are different!!

• V₁/m₁ = V₂/m₂ OR V₁/m₁ = 22.4 L/m₂

• Fewer moles = smaller volume

• More moles = larger volume

Kinetic-Molecular Theory

1. Gas particles are tiny, so their size is negligible compared to the average distance between them

2. Particles move in straight line paths, random directions, and at various speeds

3. Gas particles collide frequently with the sides of the containers and less frequently with each other and these collisions are elastic so no kinetic energy is gained or lost

4. Gas particles do not attract or repel each other because there are no IMFS

5. ↑ KE = ↑ T

Gases will deviate from ideal behaviors at _ temperatures and _ pressure, or if IMFs are _.

low, high

At these conditions, the molecules are close together, and the space between them is no longer negligible.

Gases behave most ideally at _ temperatures and _ pressure, or if IMFs are _.

high, low

At these conditions, the molecules are far apart, so they behave ideally.

Gases at different temperatures

↑ T = ↑ average particle speed, ↑ total kinetic energy, and the bell flattens out and widens

Different gases at the same temperature

• all gases are at the same temperature, so their kinetic energy is the same

• ↑ molar mass = ↓ speed

Grams Law

• if two gases are at the same temperature, the gas with the greater molar mass will be slower

• KE = ½ (mass) (velocity)²

Diffusion

gases mix together

Effusion

gases escape through a small hole

Daltons Law of Parical Pressure

• If a solution of two or more gases do not react chemically, the total pressure can be found by adding up the partial pressures

• This can only be done if the pressures you’re adding up were collected at the same conditions

• Multiple gases : P_total = P_a + P_b + P_c …

• One gas : P_A + P_total ∙ moles A/total moles (mole- fraction)

• Mole fraction = pressure fraction

Collection of Gas Over Water

• P_gas = P_total - P_water

• if pressures are in different units, convert them!

• this only works with non-polar gases at won’t dissolve!

• temp of gas =temp of room

• If we move the collection tube up/down so that the water inside the tube is equal to the level of water in the beaker, we can assume the total pressure of the gas and water vapor together us the same as the atmospheric pressure

• One the collection tube has been moved into place, we can record the volume of gas collected

• Then you can calculate the moles!

Solubility Rules for Ionic Substances

S - odium (group 1)

N - itrate (NO₃^-)

A - mmonium (NH₄⁺)

P - otasium (group 1)

AB_(s) → A_(aq)⁺ + B_(aq)^-

Hydration

process where the ends of water will interact with the cations and anions in an ionic compound

Dissociation

when an ionic compound dissolves into ions

Solubility Rules for Covalent Substances

• “Like dissolves like”

• AB → AB_(aq)

Electrolytes

substances that conduct electricity and are made of dissolves ions

Strong Electrolytes

• soluble ionic compounds

• strong acids (HI, HBr, HCl, HNO₃, H₂SO₄, HClO₃, HClO₄)

• strong bases (group 1 and 2 Hydroxides except for Be and Mg)

Weak Electrolytes

• Weak acids

• Weak bases

Non-Electrolytes

• non-soluble ionic compounds

• all covalent compounds

Separating Gases

Gases will always dissolve in each other because all IMFs have been broken

Interparticle interactions when forming solutions

1. Solute-solute interactions must be broken in an endothermic process

2. Solvent-solvent interactions must be broken in an endothermic process

3. Solvent-solute interactions must be formed in an exothermic process

If the new interactions that are formed are comparable (similar) in strength to the original interactions…

then the two substances are more likely to dissolve in each other (soluble/miscible)

If the new interactions that are formed are much weaker (different) than the original interactions…

then the two substances are more likely to not dissolve in each other (insoluble/immiscible)

Substances with _ intermolecular interactions tend to be miscible (soluble) in one another.

similar

The _ the newly formed solute-solvent interactions are, the _ the amount of solubility will be.

stronger, greater

Ionic

1. Dominate attraction

2. Strength

1. Ionic bond

2. Strong

Polar Covalent

1. Dominate attraction

2. Strength

3. Solubility with same type

1. Dipole-Dipole / HBF

2. Kinda strong

3. soluble

Non-Polar Covalent

1. Dominate attraction

2. Strength

3. Solubility with same type

1. London Dispersion Forces

2. Weak

3. soluble

Ionic + Non-Polar Covalent

1. New attraction

2. Strength

3. Solubility

1. Ion Induced Dipole

2. Weak

3. Insoluble

Ionic + Polar Covalent

1. New attraction

2. Strength

3. Solubility

1. Ion-Dipole

2. Kinda Strong

3. Often Soluble

   Depends on how strong the ionic bond is (↑ Lattice Energy = more difficult to break and less soluble)

Polar Covalent + Non-Polar Covalent

1. New attraction

2. Strength

3. Solubility

1. Dipole Induced Dipole

2. Weak

3. Insoluble

   Weak new attractions won’t be able to break the original kinda Strong dipole attraction

Chromatography

method of separating mixtures based on the intermolecular interactions of the substance being separated

Stationary Phase

solid which allows some substances to pass through and holds others back

Mobile Phase

gas or liquid which moves through the stationary phase, carrying the substances to be separated

_ attraction to the stationary phase and _ attraction to the mobile phase - will travel further up the chromatogram

Less, more

(component will be closer to the solvent front line)

_ attraction to the stationary phase and _ attraction to the mobile phase - will not travel as far up the chromatogram

More, less

(component closer to the origin line)

A larger R_f value =

less attraction to stationary phase and more attraction to the mobile phase (indicates similar attractions to identify substances)

Distillation

method of separating mixtures based on their different boiling points

A substance with a lower boiling point / boils off first has

weaker intermolecular forces

Molarity

The number of moles of solute per L of solution

M = n/L

an increase in molarity means

the substance is more concentrated

a decrease in molarity means

the substance is more diluted

Dilution Formula

V₁M₁ = V₂M₂

Spectroscopy

study of matters interactions with electromagnetic radiation

Wavelength

1. Definition

2. symbol

3. units/constants

1. distance of one completely wave cycle

2. λ

3. m, nm, cm, km

Frequency

1. Definition

2. symbol

3. units/constants

1. number of wave cycles that pass a point in one second

2. ƒ

3. Hz, S^-1, /s

Speed of Light

1. Definition

2. symbol

3. units/constants

1. constant at which all forms of light travel

2. C

3. 2.998 × 10⁸ m/s

Energy

1. symbol

2. units/constants

1. E

2. Joules/KJ

Planck’s Constant

1. symbol

2. units/constants

1. h

2. 6.626 × 10^-34 J ∙ S

Light Formulas

C = λf

E = hf

E = hC/λ

Electromagnetic Spectrum

1. Radio Waves (Red) longest wavelength, lowest frequency and energy

2. Microwaves (rotation)

3. Infrared (vibrations)

4. Visible Light (changes in electron energy levels)

5. Ultra-Violet (changes in electron energy levels)

6. X-Rays (changes in electron structure)

7. Gamma Rays (Purple) shortest wavelength, highest frequency and energy